Transition metals - complex ions -5

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  • Created by: Shannon
  • Created on: 23-04-15 08:55
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  • Transition metals - complex ions, ligand reactions & copper and chromium
    • Complex ion formation
      • A complex ion is a metal ion surrounded by datively covalently bonded ligands
        • Ligands donate a pair of electrons to a central metal ion
          • Therefore must have more than one lone pairs of electrons
            • One lone pair = monodentate ligand
            • Two lone pairs = bidentate ligand
            • More than two lone pairs = polydentate
      • The coordination number is the number of dative covalent bonds that are formed around the central metal ion
        • Usually 4 or 6, but can be 2
        • 6 dative covalent bonds = octahedral shape
        • 4 dative covalent bonds = tetrahedral shape
        • 2 dative covalent bonds = linear shape
        • [Pt(NH3)2Cl2] exists as a square planar complex
      • Ligands can be swapped in ligand subsitution reactions
        • This generally causes a colour change
      • Transition metal complexes are identified by their colour
    • Ligand reactions
      • If ligands are the same size, the coordination and shape of the complex ion does not change
        • Same
          • H2O and NH3
          • [Cr(H2O)6]3+ + 6NH3 -> [Cr(NH3)6]3+ + 6H2O
            • Colour change: violet -> purple
        • Different
          • H2O and Cl-
          • [Cu(H20)6]2+ + 4Cl- -> [CuCl4]2- + 6H2O
            • Colour change: blue -> yellow
      • Generally, ligand reactions can be reversed, but if the new ligands forms stronger bonds the change is harder to reverse
        • Monodentate < Bidentate < Polydentate
        • More stable
      • Adding NaOH or NH3 to metal aqua-ions produces insoluble metal hydroxides
        • With OH- or NH3
          • [Cu(H2O)6]2+ -> Cu(H2O)4(OH)2
            • Blue solution to blue precipitate
          • [Fe(H2O)6]2+ -> Fe(H2O)4(Oh)2
            • Green solution to green precipitate
          • [Fe(H2O)6]3+ -> Fe(H2O)3(OH)3
            • Yellow solution to brown precipitate
          • [Cr(H2O)6]3+ -> Cr(H2O)3(OH)3
            • Violet solution to green precipitate
          • [Mn(H2O)6]2+ -> Mn(H2O)4(OH)2
            • Pale pink solution to brown precipitate
          • [Ni(H2O)6]2+ -> Ni(H2O)4(OH)2
            • Green solution to green precipitate
          • [Zn(H2O6)]2+ -> Zn (H2O)4(OH)2
            • Colourless solution to white precipitate
        • With excess OH-
          • Cu = no change
          • Fe 2+ = no change
          • Fe 3+ = no change
          • Cr = green solution
          • Mn = no change
          • Ni = no change
          • Zn = colourless solution
        • With excess NH3
          • Cu = deep blue solution of [Cu(NH3)4(H2O)2]2+
          • Fe 2+ = no change
          • Fe 3+ = no change
          • Cr = purple solution of [Cr(NH3)6]3+
          • Mn = no change
          • Ni = blue solution of [Ni(NH3)6]2+
          • Zn = colourless solution of [Zn(NH3)4]2+
    • Copper and chromium
      • Copper ions usually exist in oxidation states of 1+ or 2+
        • Copper(I) ions are unstable in aqueous solution and disproportionate to giver copper and copper(II) ions
          • If 'suitable' ligands are present, Copper(I) ions are stabilised
        • Cu2+ ions are reduced to Cu if a more electropositive metal is present
          • Displacement reaction
      • Chromium ions usually exist in oxidation states of 2+, 3+ or 6+
        • Because chromium has many oxidation states it can take part in many redox reactions
          • Cr2O72-
            • Cr in 6+ ocxidation state
            • Orange
          • CrO42-
            • Cr oxidation state of 6+
            • Yellow
          • Cr3+
            • Green
          • Cr2+
            • Blue
        • Chromium hydroxide is amphoteric
          • Can act as an acid and a base

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